Hydrogen (H2) is currently the leading candidate for a fuel to replace gasoline/diesel fuel in powering the nation's transportation fleet. There are a number of difficulties and technological barriers associated with hydrogen that must be solved in order to realize this “hydrogen economy”. Inadequate storage systems for on-board transportation hydrogen are recognized as a major technological barrier (see, for example, “The Hydrogen Economy: Opportunities, Costs, Barriers, and R&D Needs,” National Academy of Engineering (NAE), Board on Energy and Environmental Systems, National Academy Press (2004).
One of the general schemes for storing hydrogen relates to using a chemical compound or system that undergoes a chemical reaction to evolve hydrogen as a reaction product. In principle, this chemical storage system is attractive, but systems that have been developed to date involve either: (a) hydrolysis of high-energy inorganic compounds where the evolution of hydrogen is very exothermic (sodium borohydride/water as in the Millennium Cell's HYDROGEN ON DEMAND®, and lithium (or magnesium) hydride as in SAFE HYDROGEN®, for example), thus making the cost of preparing the inorganic compound(s) high and life-cycle efficiency low; or (b) dehydrogenation of inorganic hydride materials (such as Na3AlH6/NaAlH4, for example) that release hydrogen when warmed but that typically have inadequate mass storage capacity and inadequate refueling rates.
Inorganic compounds referred to in (a), above, produce hydrogen according to the chemical reactionMHx+X H2O→M(OH)x+X H2  (1)where MHx is a metal hydride, and M(OH)x is a metal hydroxide. This reaction is irreversible.
Inorganic hydride materials referred to in (b), above, produce hydrogen according to the following chemical reaction, which is reversible with H2 (hydrogen gas):MHx=M+x/2H2  (2)where MHx is a metal hydride, M is metal and H2 is hydrogen gas. By contrast to the first reaction, which is irreversible with H2, the second reaction is reversible with H2.
A practical chemical system that evolves hydrogen yet does not suffer the aforementioned inadequacies would be important to the planned transportation sector of the hydrogen economy. This same practical chemical system would also be extremely valuable for non-transportation H2 fuel cell systems, such as those employed in laptop computers and other portable electronic devices, and in small mechanical devices such as lawnmowers where current technology causes significant pollution concerns.
Any heat that must be input to evolve the hydrogen represents an energy loss at the point of use, and any heat that is evolved along with the hydrogen represents an energy loss where the chemical storage medium is regenerated. Either way, energy is lost, which diminishes the life-cycle efficiency. For most organic compounds, such as in those shown in equations 3-5 below, hydrogen evolution reactions are very endothermic, and the compounds are incompetent to evolve hydrogen at ambient temperature (i.e. thermodynamically incapable of evolving H2 at significant pressure at ambient temperature). For temperatures less than about 250-400 degrees Celsius, the equilibrium pressure of hydrogen over most organic compounds is very small. As a consequence, most common organic compounds require heating above about 250 degrees Celsius, and the continual input of high-grade heat to maintain this temperature, in order to evolve hydrogen at a useful pressure.CH4→C+2H2 ΔH0=+18 kcal/mol ΔG0=+12 kcal/mol  (3)6CH4→cyclohexane+6H2 ΔH0=+69 kcal/mol ΔG0=+78 kcal/mol  (4)cyclohexane→benzene+3H2 ΔH0=+49 kcal/mol ΔG0=+23 kcal/mol  (5)
Most organic compounds have hydrogen evolution reactions that are endergonic (i.e. having a net positive standard free energy of reaction change, i.e. ΔG0>0) and their ambient temperature equilibrium hydrogen pressure is very low, practically unobservable. Thus, most organic compounds are unsuitable for hydrogen storage, based on both life-cycle energy efficiency and delivery pressure considerations. Decalin, for example, evolves hydrogen to form naphthalene when heated to about 250 degrees Celsius in the presence of a catalyst (see, for example, Hodoshima et al. in “Catalytic Decalin Dehydrogenation/Naphthalene Hydrogenation Pair as a Hydrogen Source for Fuel-Cell Vehicle,” Int. J. Hydrogen Energy (2003) vol. 28, pp. 1255-1262, incorporated by reference herein). Hodoshima et al. use a superheated “thin film” reactor that operates at a temperature of at least 280 degrees Celsius to produce hydrogen from decalin at an adequate rate and pressure. Thus, this endothermic hydrogen evolution reaction requires both a complex apparatus and high-grade heat, which diminishes the life-cycle energy efficiency for hydrogen storage.
Boranes have high hydrogen storage capacities and have attracted interest for use as hydrogen storage materials for transportation, but the difficulty of manufacturing borane compounds, and the life-cycle energy inefficiency of the chemical process presently used for their manufacture, has prevented their widespread use.
Owing to its commercial availability, NaBH4 (sodium borohydride) is a starting material typically used to prepare borane compounds. Diborane (B2H6), for example, is prepared by reacting NaBH4 with BF3. Borohydride compounds (i.e. compounds containing the BH4 anion or other anionic B—H groups) are generally prepared by reacting alkoxyborates with active metal hydrides e.g. NaH or NaAlH4. Sodium borohydride itself (NaBH4), for example, is commercially prepared using the known Schlessinger process, which involves reacting sodium hydride (NaH) with trimethoxyboron (B(OCH3)3). While convenient to practice on a small or intermediate laboratory or commercial scale, these reactions are not energy-efficient; the reaction of NaH with B(OCH3)3 is exothermic, and NaH is itself formed in the exothermic reaction of Na metal with H2, so overall, about 22 kcal of heat are released per B—H bond that is formed.
Other means are known for forming B2H6. The best known is the reaction of BCl3 with H2 at high temperature to make BHCl2 and HCl. Significant equilibrium conversion is possible only if the temperature is on the order of about 600 degrees Celsius or more, and the product mixture must be rapidly quenched, typically within a few seconds, to a temperature below about 100 degrees Celsius to allow BHCl2 to disproportionate to B2H6 and BCl3. The quenched mixture must be separated rapidly before the B2H6 back-reacts with the HCl coproduct. BCl3 and HCl are both highly corrosive. Their corrosive properties in combination with the difficulties of heat management make this process costly to practice.
Presently, there is no energy efficient means available for preparing boranes.
Methods and systems that employ chemical compounds for storing and evolving hydrogen at ambient temperature with minimal heat input remain highly desirable.